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Home , Partial pressure

Gases

1 Chapter 5

y Substances That Exist as (5.1) y of a (5.2) y The (5.3) y The Equation (5.4) y Dalton’s Law of Partial Pressure (5.5) y The Kinetic Molecular Theory of Gases (5.6) y Deviation from Ideal Behavior (5.7)

2 y Single ideal gases – (5.1-5.4) ◦ Properties of a gas (5.1, 5.2) ◦ Experimental behaviors of gases (5.3) ◦ (5.4) x STP x Calculations involving the ideal gas law ◦ Combined gas law (5.3-5.4) ◦ Stoichiometry (5.4) y of gases – (5.5) ◦ Dalton’s Law of Partial ◦ Wetting of a gas y Kinetic Molecular Theory – (5.6) ◦ Energy ◦ Molecular speed (rms) ◦ and effusion y Real gases – (5.7) ◦ 3 5.1 Substances That Exist as Gases Representation of matter:

Macroscopic versus particle

What are four properties of gases?

4 5.1 Substances That Exist as Gases

• Gases assume the and shape of their containers. • Gases are the most compressible state of matter. • Gases will mix evenly and completely when confined to the same container. • Gases have much lower densities than and .

Margin Figure, p. 138 5 5.1 Substances That Exist as Gases

Margin Figure, p. 138 6 5.1 Substances That Exist as Gases

p. 138

7 5.1 Substances That Exist as Gases y Gases do not interact with other molecules (if they do collide, it is elastic)

y Gas molecules themselves have no volume (not true but consider the vast space between each ) Ideal Gases Ideal y There is no attraction between molecules (all gases are the same)

8 5.2 Pressure of a Gas Key Questions: y What is pressure? y How is pressure measured? y What is pressure? ◦ On a particle level ◦ On a macroscopic level

9 5.2 Pressure of a Gas

Figure 5.2, p. 140 Figure 5.3, p. 141 10 5.2 Pressure of a Gas Key Questions: y What is pressure? y How is pressure measured? y What is pressure? ◦ On a particle level ◦ On a macroscopic level y What are the standard units for pressure?

11 5.3 The Gas Laws Key Questions: y How do we represent a proportionality? y How do we convert this to an equality? y What is the relationship between ◦ Pressure and volume ◦ and volume ◦ Amount and volume y How do these combine?

12 5.3 The Gas Laws

Figure 5.5, p. 144 Boyle’s Law Law Boyle’s Pressure and Volume and Pressure

Figure 5.6, p. 145 13 5.3 The Gas Laws

Figure 5.5, p. 144 Charles’s Law Law Charles’s Temperature and Volume and Temperature

Figure 5.8, p. 145 14 5.3 The Gas Laws

Figure 5.5, p. 144 Charles’s Law Law Charles’s Temperature and Pressure and Temperature

Figure 5.8, p. 145 15 5.3 The Gas Laws

Figure 5.5, p. 144 Avogadro’s Law Avogadro’s Amount and Volume Amount and

Figure 5.9, p. 147 16 5.4 The Ideal Gas Equation 5.4 The Ideal Gas Equation Key Questions: y What is “R”? y What are the units on R y What is STP? y What is the value of R

Figure 5.10, p. 148

18 5.4 The Ideal Gas Equation Key Questions: y When can a gas be assumed to be ideal (when can we use this equation)? y What can be calculated directly from the ideal gas law? y What can be calculated indirectly from the ideal gas law? y How can the properties of a gas (P, V, T) be useful in a reaction?

19 Gases and Stoichiometry When 3.1 atm of reacts with 2.7 atm of at a constant temperature in a closed rigid container, what is the theoretical yield (in atm) of ?

Figure 5.12, p. 152

20 Chapter 5 Gases – Practice What is the pressure (in atm) of 24.0 g of in a 25.0 L container at 25 oC? A. 1.7 atm B. 0.84 atm C. 0.14 atm D. 0.070 atm

What is the molar mass of an ideal gas which has a density of 0.901 g·L–1 at STP? Chapter 5 –Chapter Practice A. 24.5 g·mol–1 B. 22.4 g·mol–1 C. 20.2 g·mol–1 D. 0.901 g·mol–1

21 Chapter 5 Gases – Practice Chapter 5 –Chapter Practice

22 5.5 Dalton’s Law of Partial Pressures

Key Definitions: y partial pressure ◦ The pressures of individual gas components in a mixture y Dalton’s law of partial pressures ◦ Total pressure of a mixture of gases is just the sum of the pressures that each gas would exert if it were present alone y fraction ◦ Dimensionless quantity that expresses the ratio of the number of moles of one component to the number of moles of all components present 23 Dalton’s Law of Partial Pressures

Figure 5.13, p. 154

24 The reaction of oxygen and to form water is utilized in fuel cells. In the figure below the volume of the rigid container is 10.0 L, the temperature is constant at 55oC and each symbol represents 0.050 mol of gas. Water molecules are not explicitly shown in any phase.

If the total pressure is 1243 mmHg, what are the partial pressures of all gases?

If the reaction continues to completion, what is the resulting partial pressure of oxygen and hydrogen?

25 5.5 Dalton’s Law of Partial Pressures

Key Questions: y Does partial pressure change with container size? ◦ with temperature? y What is wetting of a gas? ◦ Why is this a useful experimental technique? ◦ Why does Dalton’s Law of Partial Pressures apply?

Table 5.2, p. 157 26 5.5 Dalton’s Law of Partial Pressures

Figure 5.14, p. 157 27 5.5 Dalton’s Law of Partial Pressures More Practice:

Hydrogen peroxide (H2O2) decomposes into water and oxygen. This process can be sped up through the use of a catalyst.

If a 1.50 mL sample of a hydrogen peroxide is decomposed with a catalyst and 79.0 mL of gas is collected over water at 22.5ºC and under 1.015 atm of pressure, what is the of hydrogen peroxide (in water solution) (%v/v)?

The pressure of water at 22.5ºC is 20.20 mmHg and the density of hydrogen peroxide is 1.46 g·mL–1.

28 Hydrogen peroxide (H2O2) decomposes into water and oxygen. This process can be sped up through the use of a catalyst. If a 1.50 mL sample of a hydrogen peroxide solution is decomposed with a catalyst and 79.0 mL of gas is collected over water at 22.5ºC and under 1.015 atm of pressure, what is the concentration of hydrogen peroxide (in water solution) (%v/v)? The of water at 22.5ºC is 20.20 mmHg and the density of hydrogen peroxide is 1.46 g·mL–1. 1. What is the total pressure in mmHg? 2. What is the partial pressure of oxygen? 3. How many moles of oxygen were produced? 4. What is the balanced chemical equation? 5. How many moles of peroxide were present? 6. What volume of peroxide was present? 7. What is the %v/v concentration?

29 5.5 Dalton’s Law of Partial Pressures More Practice:

The reaction N2O4(g) Æ 2NO2(g) does not actually go to completion at all . If 35.0 g of N2O4 is originally placed in a 2.0 L container at 100°C and allowed to react, the final pressure in the container is 10.6 atm. What is the partial pressure of NO2 and N2O4 at equilibrium and what is the percent yield of the reaction under these conditions?

30 The reaction N2O4(g) Æ 2NO2(g) does not actually go to completion at all temperatures. If 35.0 g of N2O4 is originally placed in a 2.0 L container at 100°C and allowed to react, the final pressure in the container is 10.6 atm. What is the partial pressure of NO2 and N2O4 at equilibrium and what is the percent yield of the reaction under these conditions?

1. How many moles of N2O4 are present initially? 2. How many total moles of gases were present after the reaction ended? 3. How can the completion of the reaction be modeled using a variable (“x”) and the correct mole ratios (remember, the reactants are consumed (or decreasing number of moles) and the products are formed (or increasing number of moles)?

4. How many moles of NO2 were experimentally formed? 5. What is the percent yield? 6. What are the partial pressures of both gases?

31 5.6 The Kinetic Molecular Theory of Gases Key Questions: y (Review) What do we assume when we say a gas is “ideal”? y How are average kinetic energy and temperature related for an ideal gas? ◦ Will all ideal gases, regardless of type have the same average kinetic energy at the same temperature? y How are molecular speed and temperature related for an ideal gas?

32 5.6 The Kinetic Molecular Theory of Gases

Figure 5.15, p. 161 33 5.6 The Kinetic Molecular Theory of Gases

34 5.6 The Kinetic Molecular Theory of Gases

35 5.6 The Kinetic Molecular Theory of Gases

36 5.6 The Kinetic Molecular Theory of Gases

37 5.6 The Kinetic Molecular Theory of Gases

38 5.6 The Kinetic Molecular Theory of Gases

39 5.6 The Kinetic Molecular Theory of Gases

40 5.6 The Kinetic Molecular Theory of Gases

41 5.6 The Kinetic Molecular Theory of Gases

42 5.6 The Kinetic Molecular Theory of Gases Key Questions: y What do we assume when we say a gas is “ideal”? y How are average kinetic energy and temperature related for an ideal gas? ◦ Will all ideal gases, regardless of type have the same average kinetic energy at the same temperature? y How are molecular speed, temperature and gas identity related for an ideal gas?

43 Will the speed of all gases be the same at one temperature?

Figure 5.15, p. 161 44 Diffusion/Effusion Key Definition: y Diffusion ◦ Gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic properties

Key Question: y Why does a gas, traveling so fast, take such a long time to diffuse?

45 Figure 5.16, p. 163 46 Diffusion/Effusion Key Definition: y Effusion ◦ The process by which a gas under pressure escapes from one compartment of a container to another by passing through a small opening.

Figure 5.18, p. 164

47 5.7 Deviation from Ideal Behavior

y Under what conditions can a gas be assumed to be ideal?

Figure 5.20, p. 165 and Table 5.3, p. 166 48 5.7 Deviation from Ideal Behavior

y Under what conditions can a gas be assumed to be ideal?

Figure 5.19, p. 165

49 Chapter 5 Gases – Practice

Which ideal gas will have the greatest average kinetic energy at STP?

A. H2 B. He C. Ne D. All ideal gases have the same average kinetic energy for a given temperature.

What is the root mean square speed of fluorine at room temperature (25 oC)? A. less than 50 m·s–1

Chapter 5 –Chapter Practice B. between 50 m·s–1 and 275 m·s–1 C. between 275 m·s–1 and 500 m·s–1 D. greater than 500 m·s–1

50 Chapter 5 Gases – Practice Chapter 5 –Chapter Practice

51